Acids, Bases, Salts, pH, and Buffers: Chemistry Foundations for Anatomy & Physiology

Chemistry Part II: Acids, Bases, Salts, pH and Buffers

Dissociation

Dissociation is a fundamental chemical process in which molecules, such as acids, bases, or salts, separate into smaller particles like atoms or ions. This process is essential for understanding how substances interact in biological systems.

• Dissociation: The splitting of a compound into ions when dissolved in water.

• When an acid dissolves in water, it breaks a covalent bond, releasing a proton (H+) and a negative ion.

• Example: Water can dissociate into hydrogen ions (H+) and hydroxide ions (OH-).

Equation:

$\mathrm{H_2O \rightleftharpoons H^+ + OH^-}$

In pure water, only a small fraction of molecules dissociate at any time.

Acids

Acids are substances that release hydrogen ions (H+) when they dissociate in water. The strength of an acid depends on its ability to dissociate completely or partially.

• Strong acids (e.g., HCl) dissociate completely in solution:

• $\mathrm{HCl \rightarrow H^+ + Cl^-}$

• Weak acids (e.g., acetic acid, carbonic acid) only partially dissociate:

• $\mathrm{CH_3COOH \rightleftharpoons CH_3COO^- + H^+}$

• $\mathrm{H_2CO_3 \rightleftharpoons HCO_3^- + H^+}$

• Acids are H+ donors. Acidic solutions have high H+ concentrations.

Example: Hydrochloric acid (HCl) is a strong acid, while acetic acid (CH3COOH) is a weak acid.

Strong Versus Weak Acids

The distinction between strong and weak acids is based on the degree of dissociation in water.

• Strong acids dissociate completely, releasing all possible H+ ions.

• Weak acids dissociate partially, releasing fewer H+ ions and having a smaller effect on pH.

• Double-headed arrows ($\rightleftharpoons$) indicate reversible reactions typical of weak acids.

Comparison Table:

Type	Dissociation	Effect on pH
Strong Acid	Complete	Large decrease
Weak Acid	Partial	Small decrease

Bases

Bases are substances that release hydroxide ions (OH-) in water or accept hydrogen ions (H+). Their strength depends on their ability to dissociate and bind H+ ions.

• Strong bases (e.g., NaOH) dissociate completely:

• $\mathrm{NaOH \rightarrow Na^+ + OH^-}$

• OH- ions bind to H+, removing them from solution.

• Weak bases (e.g., NH3) dissociate partially:

• $\mathrm{NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-}$

• Bases are H+ acceptors. Basic solutions have low H+ concentrations.

Example: Sodium hydroxide (NaOH) is a strong base, while ammonia (NH3) is a weak base.

Strong Versus Weak Bases

• Strong bases dissociate completely and readily bind free H+.

• Weak bases bind fewer H+ ions and have a smaller effect on pH.

Comparison Table:

Type	Dissociation	Effect on pH
Strong Base	Complete	Large increase
Weak Base	Partial	Small increase

Bases and Salts

Salts are ionic compounds formed from the reaction of acids and bases. They consist of any cation except H+ and any anion except OH-.

• Acids and bases can neutralize each other to form a salt and water:

• $\mathrm{Acid + Base \rightarrow Salt + Water}$

• Example: $\mathrm{HCl + NaOH \rightarrow NaCl + H_2O}$

• Salts dissociate in water to produce ions (e.g., Na+, Cl-).

• Salts are important electrolytes in body fluids, helping regulate nerve and muscle function.

Electrolytes: Inorganic compounds that dissociate into ions in solution, such as Na+, K+, Ca2+, Cl-, HCO3-.

pH Scale

The pH scale measures the concentration of hydrogen ions (H+) in a solution, indicating its acidity or alkalinity.

• pH is defined as:

$\mathrm{pH = -\log [H^+]}$

• pH ranges from 0 (most acidic) to 14 (most basic).

• pH = 7 is neutral (equal H+ and OH-).

• pH < 7 is acidic (more H+ than OH-).

• pH > 7 is basic (more OH- than H+).

• The pH scale is logarithmic: a change of 1 pH unit equals a tenfold change in [H+].

Example: A solution with pH 5 has 10 times more H+ than a solution with pH 6.

Buffers and pH Regulation

Buffers are mixtures of weak acids and their corresponding bases that minimize changes in pH when acids or bases are added to a solution. They are vital for maintaining homeostasis in living organisms.

• Buffers work by reversibly binding H+ ions.

• If H+ increases, the buffer's base binds H+ to lower pH.

• If H+ decreases, the buffer's acid releases H+ to raise pH.

• The carbonic acid/bicarbonate buffer system is crucial in blood:

$\mathrm{H_2CO_3 \rightleftharpoons HCO_3^- + H^+}$

• When blood pH rises (becomes more basic), carbonic acid dissociates to release H+.

• When blood pH falls (becomes more acidic), bicarbonate binds H+ to form carbonic acid.

Example: If blood pH goes above normal, the buffer system shifts to remove excess OH- and restore pH balance.

Summary Table: Acid/Base Properties

Property	Acid	Base
Ion Released	H+	OH-
pH Effect	Decreases	Increases
Strength	Strong/Weak	Strong/Weak
Example	HCl, CH3COOH	NaOH, NH3

Additional info: Electrolyte balance is essential for nerve impulse transmission, muscle contraction, and maintaining fluid balance in the body. The carbonic acid/bicarbonate buffer system is the primary mechanism for regulating blood pH in humans.